By Chris Deziel | Updated March 24, 2022
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The mass number of an element is shown directly beneath its symbol in the periodic table. It’s expressed in atomic mass units (amu), which equal grams per mole.
Each element is defined by its atomic number—the count of positively charged protons in its nucleus. Hydrogen, for instance, has one proton; oxygen has eight. The periodic table arranges elements by increasing atomic number.
The mass number, however, is the total count of protons plus neutrons in the nucleus. Since neutrons carry roughly the same mass as protons but no charge, they must be included when calculating the element’s mass. For example, the most common isotope of oxygen has eight protons and eight neutrons, giving it a mass number of 16.
Look beneath the symbol of any element in a reputable periodic table—this is the mass number. You may notice a decimal, which is actually the relative atomic mass, the weighted average of all naturally occurring isotopes. Because most elements have multiple isotopes, the average is rarely an integer.
For instance, the periodic table lists hydrogen’s mass as 1.008, carbon’s as 12.011, and oxygen’s as 15.99. Uranium (atomic number 92) has three natural isotopes, giving it a relative mass of 238.029. In routine calculations, scientists round to the nearest whole number.
Since the early 20th century, the atomic mass unit (amu) has been the standard. It is defined as exactly one‑twelfth the mass of an isolated carbon‑12 atom. Consequently, 1 amu equals 1 gram per mole. Therefore, one mole of hydrogen weighs 1 gram, one mole of carbon 12 grams, and one mole of uranium 238 grams.
These definitions, adopted by IUPAC and used by NIST, ensure consistency across scientific literature and chemical calculations.
Reference: IUPAC Technical Report, 2021; NIST Chemistry WebBook.
